THE PERIODIC TABLE
The periodic table is a list of all known elements arranged in order of increasing atomic number, from 1 to 106. In addition to this, the elements are arranged in such a way that atoms with the same number of shells are placed together, and atoms with similar electronic configurations in the outer shell are also placed together. This is achieved as follows:
The elements are arranged in rows and columns. Elements with one shell are placed in the first row (ie H and He), Elements with two shells are placed in the second row (Li to Ne) and so on.
A row of elements thus arranged is called a period.
In addition, the elements are aligned vertically (in columns) with other elements in different rows, if they share a similar outer-shell electronic configuration. For example, elements with outer-shell configuration ns1 are all placed in the same column ( Li, Na, K, Rb, Cs, Fr).
A column of elements thus arranged is called a group.
According to these principles, the periodic table can be constructed as follows:
I II III IV V VI VII 0
|Cs||Ba||La||Ce – Lu||Hf||Ta||W||Re||Os||Ir||Pt||Au||Hg||Tl||Pb||Bi||Po||At||Rn|
|Fr||Ra||Ac||Th – Lw|
Since the electronic configurations of H and He are unusual, they do not fit comfortably into any group. They are thus allocated a group based on similarities in physical and chemical properties with other members of the group.
He is placed in group 0 on this basis, but hydrogen does not behave like any other element and so is placed in a group of its own.
The elements Ce – Lu and Th – Lw belong in the periodic table as shown above. However if they are placed there periods 6 and 7 do not fit onto a page of A4, so they are placed below the other elements in most tables.
All elements belong to one of four main blocks: the s-block, the p-block, the d-block and the f-block.
The s-block elements are all those with only s electrons in the outer shell.
The p-block elements are all those with at least one p-electron in the outer shell.
The d-block elements are all those with at least one d-electron and at least one s-electron but no f or p electrons in the outer shell.
The f-block elements are all those with at least one f-electron and at least one s-electron but no d or p electrons in the outer shell.
The physical and chemical properties of elements in the Periodic Table show clear patterns related to the position of each element in the Periodic Table. Elements in the same group show similar properties, and properties change gradually on crossing a Period.
As atomic number increases, the properties of the elements show trends which repeat themselves in each Period of the Periodic Table. These trends are known as Periodic Trends and the study of these trends in known as Periodicity.
- First ionisation energy
|The first ionisation energy of an element is the energy required to remove one electron from each of a mole of free gaseous atoms of that element.|
It can also be described as the energy change per mole for the process:
M(g) —> M+(g) + e
The amount of energy required to remove an electron from an atom depends on the number of protons in the nucleus of the atom and on the electronic configuration of that atom.
The first ionisation energies of the first 20 elements in the periodic table is shown below:
There are various trends in this graph which can be explained by reference to the proton number and electronic configuration of the various elements. A number of factors must be considered:
– Energy is required to remove electrons from atoms in order to overcome their attraction to the nucleus. The greater the number of protons, the greater the attraction of the electrons to the nucleus and the harder it is to remove the electrons. The number of protons in the nucleus is known as the nuclear charge.
– The effect of this nuclear charge, however, is cancelled out to some extent by the other electrons in the atom. Each inner shell and inner sub-shell electron effectively cancels out one unit of charge from the nucleus. This is known as shielding.
– The outermost electrons in the atom thus only feel the residual positive charge after all inner shell and inner sub-shell electrons have cancelled out much of the nuclear charge. This residual positive charge is known as the effective nuclear charge.
– Electrons repel each other, particularly when they are in the same orbital. The degree of repulsion between the outermost electrons affects the ease with which electrons can be moved.
When considering trends in ionisation energies, it is thus necessary to consider 4 factors:
- nuclear charge
- effective nuclear charge
- electron repulsion
The trends in first ionisation energies amongst elements in the periodic table can be explained on the basis of variations in one of the four above factors.
Trend across period 1
Compare the first ionisation energies of H and He. Neither have inner shells, so there is no shielding. He has two protons in the nucleus; H only has one. Therefore the helium electrons are more strongly attracted to the nucleus and hence more difficult to remove.
The first ionisation energy of He is thus higher than that of H.
Since H and He are the only atoms whose outer electrons are not shielded from the nucleus, it follows that He has the highest first ionisation energy of all the elements. All elements (except H) have outer electrons which are shielded to some extent from the nucleus and thus are easier to remove.
So Helium has the highest first ionisation energy of all the elements.
Trends across period 2
Compare now the first ionisation energies of He (1s2) and Li (1s22s1). Li has an extra proton in the nucleus (3) but two inner-shell electrons. These inner-shell electrons cancel out the charge of two of the protons, reducing the effective nuclear charge on the 2s electron to +1. This is lower than the effective nuclear charge on the He 1s electrons, +2, and so the electrons are less strongly held and easier to remove.
The first ionisation energy of Li is thus lower than that of He.
Compare the first ionisation energies of Li (1s22s1) and Be (1s22s2). Be has one more proton in the nucleus than Li, and no extra inner-shell electrons, so the effective nuclear charge on Be is higher and the Be electrons are more strongly attracted to the nucleus.
The first ionisation energy of Be is thus higher than that of Li.
In general, the first ionisation energy increases across a period because the nuclear charge increases but the shielding remains the same.
Compare the first ionisation energies of Be (1s22s2) and B (1s22s22p1).B has one more proton in the nucleus than Be but there are also 2 extra inner sub-shell electrons. These cancel out the charge of two more of the protons, leaving an effective nuclear charge of only +1. This is less than Be (+2) so the electrons are less strongly attracted to the nucleus and thus less difficult to remove.
The first ionisation energy of B is thus lower than that of Be.
Ionisation energies decrease from group II to group III because in group III the electrons are removed from a p-orbital, so it is shielded by the s-electrons in the outer shell. Thus the effective nuclear charge decreases.
From B (1s22s22p1) to N (1s22s22p3) the proton number increases, but the number of electrons shielding the nuclear charge remains the same at 4. Thus the effective nuclear charge increases from B to N and the electrons become progressively harder to remove.
The first ionisation energy thus increases from B to N.
So far the concepts of effective nuclear charge and shielding have been used to explain the trend in first ionisation energies for the first 7 elements. They cannot, however, explain the fall between N and O. The electronic configurations of N and O must be considered more carefully:
1s 2s 2p
Note that in N the electron is removed from an unpaired orbital, but in O it is removed from a paired orbital. In a paired orbital, the two electrons share a confined space and so repel each other. They are therefore less stable and easier to remove. This repulsion effect outweighs the higher effective nuclear charge in O.
The first ionisation energy of O is thus lower than that of N.
First ionisation energies decrease from group V to group VI, since the electron removed from the group VI atom is paired, so there is more repulsion between the electrons and the electron is easier to remove.
The first ionisation energies increase as expected from O to Ne, due to the increase in effective nuclear charge.
The trend in first ionisation energies across period 2 can thus be summarised as follows:
- There is a general increase across the period as the nuclear charge increases and the shielding remains the same.
- There is a drop from Be to B because in B a 2p electron is being removed and the extra shielding from the 2s subshell actually causes a fall in the effective nuclear charge.
- There is also a drop from N to O because the electron in O is being removed from a paired orbital. The repulsion of the electrons in this orbital makes them less stable and easier to remove.
The same trend can also be found in Period 3 (Na – Ar). There is a general increase, but a drop between Mg and Al and also between P and S.
Trend down a group
The above graph also shows a clear decrease in first ionisation energy on descending a group. This can be explained in the following way:
On descending a group, the effective nuclear charge stays the same but the number of inner shells increases. The repulsion between these inner shells and the outer electrons makes them less stable, pushes them further from the nucleus and makes them easier to remove.
ii) Successive ionisation energies
The second ionisation energy of an atom is the energy required to remove one electron from each of a mole of free gaseous unipositive ions.
M+(g) —-> M2+(g) + e
Other ionisation energies can be defined in the same way:
The third ionisation energy of an atom is the energy required to remove one electron from each of a mole of bipositive ions.
M2+(g) —–> M3+(g) + e
The nth ionisation energy can be defined as the energy required for the process
M(n-1)+(g) —-> Mn+(g) + e
It always becomes progressively more difficult to remove successive electrons from an atom; the second ionisation energy is always greater than the first, the third always greater than the second and so on. There are two satisfactory explanations for this:
As more electrons are removed from an atom, the number of electrons remaining in the atom decreases. The repulsion between these electrons therefore decreases, while the number of protons remains the same. The remaining electrons are thus more stable and increasingly difficult to remove.
The difference in successive ionisation energies, however, varies widely and depends on the electronic configuration of the atom in question. The difference in successive ionisation energies of an atom can be predicted qualitatively by consideration of the effective nuclear charge on the electron to be removed and the shielding of that electron by the inner shell and inner sub-shell electrons.
Consider the successive ionisation energies of aluminium, 1s22s22p63s23p1:
The 1st ionisation energy is fairly low because the 3p electron is shielded by all the other electrons, and the effective nuclear charge is only +1.
The 2nd and 3rd ionisation energies are significantly higher than the 1st because 3s electrons are being removed and the effective nuclear charge on these electrons is +3.
1st: 578 kJmol-1, 2nd: 1817 kJmol-1, 3rd: 2745 kJmol-1
There is a huge jump to the 4th ionisation energy, since a 2p electron is now being removed. The shielding has fallen and the effective nuclear charge has risen to +9.
The 5th and 6th ionisation energies are also high.
4th: 11578 kJmol-1, 5th: 14831 kJmol-1, 6th: 18378 kJmol-1
There is another significant jump to the 7th ionisation energy, since an unpaired 2p electron is now being removed.
7th: 23296 kJmol-1, 8th: 27460 kJmol-1, 9th: 31862 kJmol-1
The next significant jump is between the 9th and 10th ionisation energies, since the 10th requires the removal of a 2s electron.
10th: 38458kJmol-1, 11th: 42655 kJmol-1
There is a huge jump to the12th ionisation energy, since a 1s electron is now being removed.
12th: 201276kJmol-1, 13th: 222313kJmol-1.
These ionisation energies could be plotted on a graph as follows:
Note that the largest jumps by far occur between the 3rd and 4th ionisation energies, and between the 11th and 12th ionisation energies. In practice only large jumps such as this are visible on such a graph.
The relative values of successive ionisation energies are therefore a direct indicator of the electronic configuration of the atom in question.
The trends can be summarised as follows:
- The successive ionisation energies of an atom always increase. The more electrons that are removed, the fewer the number electrons that remain. There is therefore less repulsion between the electrons in the resulting ion. The electrons are therefore more stable and harder to remove.
- By far the largest jumps between successive ionisation energies come when the electron is removed from an inner shell. This causes a large drop in shielding, a large increase in effective nuclear charge and a large increase in ionisation energy
By applying the above principles in reverse, it is also possible to predict the electronic structure of a species by analysis of the successive ionisation energy data:
ATOMIC AND IONIC SIZE
a) atomic size
On moving across the Periodic Table from left to right, the nuclear charge increases but the shielding stays the same. The attraction of the outer electrons to the nucleus thus increases and the outer electrons are pulled in closer. The size of the atoms therefore decreases on crossing a period – i.e. sodium is the largest atom in Period 3 and argon is the smallest.
On moving down the Periodic Table from top to bottom, the nuclear charge increases but the number of shells also increases, and the increase in shielding outweighs the increase in nuclear charge. The outermost electrons are therefore held less closely to the nucleus and drift further out. The size of the atoms therefore increases down a group – i.e. beryllium is the smallest atom in group 2 and radium is the largest.
b) ionic size
When you remove an electron from the outer shell, the repulsion between the remaining electrons decreases, and they are able to move closer to the nucleus. Cations are therefore always smaller than the corresponding atoms of the same element.
When you add an electron to the outer shell, the repulsion between the electrons increases, and they are pushed further away from the nucleus. Anions are therefore always larger than the corresponding atoms of the same element.